Understanding atomic structure is foundational to the study of inorganic chemistry. Every element is made up of atoms, which consist of a nucleus containing protons and neutrons, surrounded by electrons that occupy defined energy levels or orbitals. Over time, scientific models of the atom have evolved, from Dalton's solid sphere model to Thomson's plum pudding model, and then to Rutherford's nuclear model. The most accepted model today is the quantum mechanical model, which describes electron behavior in terms of probabilities and wavefunctions rather than fixed orbits. Electrons are arranged in shells and subshells, and their distribution is described by electron configurations. This configuration determines how atoms interact, bond, and behave chemically.

The periodic table is a systematic arrangement of elements based on their atomic number, electron configuration, and recurring chemical properties. Its structure reveals important periodic trends that influence the behavior of elements. For instance, atomic radius decreases across a period from left to right due to increasing nuclear charge, while it increases down a group as additional electron shells are added. Ionization energy, or the energy required to remove an electron, increases across a period and decreases down a group. Electronegativity, which describes an atom's ability to attract electrons in a bond, follows a similar trend. These periodic trends are crucial for predicting reactivity and bonding patterns.

The periodic table is divided into blocks (s, p, d, and f) based on electron configuration, and into groups and periods that classify elements with similar characteristics. The alkali metals, for example, are found in Group 1 and are highly reactive, especially with water. The noble gases in Group 18 are inert due to their full valence shells. Transition metals, located in the d-block, are known for their ability to form colored compounds and multiple oxidation states.